1. 7 GENERAL CHEMISTRY Thermochemistry Principles and Modern Applications
Chemistry 140 Fall 2002
TENTH EDITION
GENERAL CHEMISTRY
Principles and Modern Applications
PETRUCCI HERRING MADURA BISSONNETTE 7
Thermochemistry
In this chapter, you will learn how to deduce and write chemical formulas and how to use the information incorporated into chemical formulas. The chapter ends with an overview of the relationship between names and formulas—chemical nomenclature.
PHILIP DUTTON
UNIVERSITY OF WINDSOR
DEPARTMENT OF CHEMISTRY AND BIOCHEMISTRY
Slide 1 of 57
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

2. General Chemistry: Chapter 7
Chemistry 140 Fall 2002
Thermochemistry
Contents
7-1
Getting Started: Some
Terminology
7-2
Heat
7-3
Heats of Reaction
and Calorimetry
7-4
Work
7-5
The First Law of
Thermodynamics
7-6
Heats of Reaction: DU and DH
7-7
Indirect Determination
of DH: Hess’s Law
7-8
Standard Enthalpies of
Formation
Potassium reacts with water, liberating sufficient heat to ignite the hydrogen evolved. The transfer of heat between substances in chemical reactions is an important aspect of thermochemistry.
Thermochemistry is the branch of chemistry concerned with the heat effects that accompany chemical reactions. To understand the relationship between heat and chemical and physical changes, we must start with some basic definitions. We will then explore the concept of heat and the methods used to measure the transfer of energy across boundaries. Another form of energy transfer is work, and, in combination with heat, we will define the first law of thermodynamics.
Slide 2 of 57
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

3. 6-1 Getting Started: Some Terminology
Chemistry 140 Fall 2002
6-1 Getting Started: Some Terminology
Open system. The beaker of hot coffee transfers energy to the surroundings—it loses heat as it cools. Matter is also transferred in the form of water vapor.
Closed system. The flask of hot coffee transfers energy (heat) to the surroundings as it cools. Because the flask is stoppered, no water vapor escapes and no matter is transferred.
Isolated system. Hot coffee in an insulated container approximates an isolated system. No water vapor escapes, and, for a time at least, little heat is transferred to the surroundings. (Eventually, though, the coffee in the container cools to room temperature so this system is not completely isolated.)
FIGURE 7-1
Systems and their surroundings
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

4. General Chemistry: Chapter 7
Chemistry 140 Fall 2002
Energy, U
The capacity to do work.
Work, w
Force acting through a distance.
Kinetic Energy, eK
The energy of motion.
Potential Energy, V
The stored energy has potential to do work.
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

5. General Chemistry: Chapter 7
Chemistry 140 Fall 2002
Kinetic Energy 2 1 m
ek =
mv2
[ek ] = kg
= J 2 s
Work
w = F x d
Click 1 exposes the work formulae
Click 2 exposes the units of work, notice they are the same as the units of kinetic energy m m
= m x a x d
[w ] = kg
= J s2
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

6. General Chemistry: Chapter 7
Chemistry 140 Fall 2002
Potential Energy
Energy due to condition, position, or composition.
Associated with forces of attraction or repulsion between objects.
Energy can change from potential to kinetic.
Energy changes continuously from potential to kinetic.
Energy is lost to the surroundings.
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

7. General Chemistry: Chapter 7
7-2 Heat
Thermal Energy
Kinetic energy associated with random molecular motion.
In general proportional to temperature.
Heat and Work
q and w.
Energy changes.
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

8. General Chemistry: Chapter 7
Chemistry 140 Fall 2002
Internal energy is transferred between a system and its surroundings as a result of a temperature difference.
Heat “flows” from hotter to colder.
Temperature may change.
Phase may change (an isothermal process).
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

9. General Chemistry: Chapter 7
Calorie (cal)
The quantity of heat required to change the temperature of one gram of water by one degree Celsius.
Joule (J)
SI unit for heat
1 cal = J
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

10. General Chemistry: Chapter 7
Heat Capacity
The quantity of heat required to change the temperature of a system by one degree.
Molar heat capacity.
System is one mole of substance.
Specific heat capacity, c.
System is one gram of substance
Heat capacity
(Mass of system) x specific heat.
q =
mcT
q =
CT
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

11. Law of conservation of energy
In interactions between a system and its surroundings the total energy remains constant— energy is neither created nor destroyed.
qsystem + qsurroundings = 0
qsystem = -qsurroundings
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

12. General Chemistry: Chapter 7
TABLE 7.1 Some Specific Heat Values, J g-1 °C-1
Chemistry 140 Fall 2002
Can we explain why liquid water has a high specific heat? The answer is most certainly yes, but the explanation relies on concepts we have not yet discussed. The fact that water molecules form hydrogen bonds (which we discuss in Chapter 12) is an important part of the reason why water has a large specific heat value. Energy is stored in hydrogen bonds, particularly strong intermolecular interactions.
High Specific Heat of Compounds
Because of their greater complexity at the molecular level, compounds generally have more ways of storing internal energy than do the elements; they tend to have higher specific heats. Water, for example, has a specific heat that is more than 30 times as great as that of lead. We need a much larger quantity of heat to change the temperature of a sample of water than of an equal mass of a metal.
Lake Effect
An environmental consequence of the high specific heat of water is found in the effect of large lakes on local climates. Because a lake takes much longer to heat up in summer and cool down in winter than other types of terrain, lakeside communities tend to be cooler in summer and warmer in winter than communities more distant from the lake.
We need a much larger quantity of heat to change the temperature of a sample of water than of an equal mass of a metal.
Slide 12 of 57
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

13. 7-3 Heats of Reaction and Calorimetry
Chemistry 140 Fall 2002
7-3 Heats of Reaction and Calorimetry
Chemical energy.
Contributes to the internal energy of a system.
Heat of reaction, qrxn.
The quantity of heat exchanged between a system and its surroundings when a chemical reaction occurs within the system, at constant temperature.
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

14. Exothermic and endothermic reactions
Chemistry 140 Fall 2002
(a) An exothermic reaction. Slaked lime, Ca(OH)2 is produced by the action of water on quicklime, (CaO). The reactants are mixed at room temperature,
but the temperature of the mixture rises to 40.5°C
(b) An endothermic reaction. Ba(OH)2 • 8 H2O(s) and NH4Cl are mixed at room temperature, and the temperature falls to 5.8°C in the reaction which produces BaCl2•2 H2O and 2 NH3 and releasing 8 H2O.
Exothermic and endothermic reactions
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

15. The “Coffee-Cup” Calorimeter
Chemistry 140 Fall 2002
Determining a Heat of Reaction from Calorimetric Data
The “Coffee-Cup” Calorimeter
A simple calorimeter.
Well insulated and therefore isolated.
Measure temperature change.
qrxn = -qcal
see EXAMPLE 7-4 Determining a Heat of Reaction from Calorimetric Data
qcal is the quantity of heat producing the temperature change
in the calorimeter.
FIGURE 7-6
A Styrofoam “coffee-cup” calorimeter
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

16. General Chemistry: Chapter 7
7-4 Work
In addition to heat effects chemical reactions may also do work.
Gas formed pushes against the atmosphere.
The volume changes.
Work involved in compression or expansion of gases is called
Pressure-volume work.
FIGURE 7-7
Illustrating work (expansion) during the chemical reaction
2 KClO3(s) KCl(s) + 3 O2(g)
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

17. General Chemistry: Chapter 7
Chemistry 140 Fall 2002
w = F x d
= (m x g)
x h
(m x g)
 h
x A = A
= PV
w = -PextV
This figure is demonstrated in EXAMPLE 7-5 Calculating Pressure–Volume Work
In this hypothetical apparatus, a gas is confined by a massless piston of area A. A massless wire is attached to the piston and the gas is held back by two weights with a combined mass of 2M resting on the massless pan. The cylinder is immersed in a large water bath in order to keep the gas temperature constant. The initial state of the gas is P= 2Mg/A with a volume Vi at temperature, T.
When the external pressure on the confined gas is suddenly lowered by removing one of the weights the gas expands, pushing the piston up by the distance Dh, The increase in volume of the gas (DV) is the product of the cross-sectional area of the cylinder (A) and the distance (Dh). The final state of the gas is Pf = Mg/A, Vf, and T.
Negative sign is introduced for Pext because the system does work ON the surroundings.
When a gas expands V is positive and w is negative.
When a gas is compressed V is negative and w is positive, indicating that energy (as work) enters the system.
In many cases the external pressure is the same as the internal pressure of the system, so Pext is often just represented as P
FIGURE 7-8
Pressure-volume work
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

18. General Chemistry: Chapter 7
Chemistry 140 Fall 2002
To describe a system completely we must describe its T, P and the amount of
Substance present.
Open system. The beaker of hot coffee transfers energy to the surroundings—it loses heat as it cools. Matter is also transferred in the form of water vapor.
Closed system. The flask of hot coffee transfers energy (heat) to the surroundings as it cools. Because the flask is stoppered, no water vapor escapes and no matter is transferred.
Isolated system. Hot coffee in an insulated container approximates an isolated system. No water vapor escapes, and, for a time at least, little heat is transferred to the surroundings. (Eventually, though, the coffee in the container cools to room temperature so this system is not completely isolated.)
Systems and their surroundings
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

19. The “Coffee-Cup” Calorimeter
Chemistry 140 Fall 2002
Determining a Heat of Reaction from Calorimetric Data
The “Coffee-Cup” Calorimeter
A simple calorimeter.
Well insulated and therefore isolated.
Measure temperature change.
qrxn = -qcal
see EXAMPLE 7-4 Determining a Heat of Reaction from Calorimetric Data
qcal is the quantity of heat producing the temperature change
in the calorimeter.
FIGURE 7-6
A Styrofoam “coffee-cup” calorimeter
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

20. Some contributions to the internal energy of a system
Chemistry 140 Fall 2002
Some contributions to the internal energy of a system
Internal Energy, U.
Total energy (potential and kinetic) in a system.
Translational kinetic energy.
Molecular rotation.
Bond vibration.
Intermolecular attractions.
Chemical bonds.
Electrons.
The models represent water molecules, and the arrows represent the types of motion they can undergo. In the intermolecular attractions between water molecules, the symbols ∂+ and ∂- signify a separation of charge, producing centers of positive and negative charge that are smaller than ionic charges. These intermolecular attractions are discussed in Chapter 12.
FIGURE 7-9
Some contributions to the internal energy of a system
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

21. The First Law of Thermodynamics
A system contains only internal energy.
A system does not contain heat or work.
These two are means by which the system exchanges energy with the surrounding..
Law of Conservation of Energy
The energy of an isolated system is constant
U = q + w
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

22. The First Law of Thermodynamics
An isolated system is unable to exchange either heat or work with its surroundings, so that DUisolated system = 0, and we can say:
The energy of an isolated system is constant.
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

23. General Chemistry: Chapter 7
Any property that has a unique value for a specified state of a system is said to be a State Function.
◦ Water at K and 1.00 atm is in a specified state.
◦ d = g/mL
◦ This density is a unique function of the state.
◦ It does not matter how the state was established.
The value of a function of state depends on the state of the system, and not on how that state was established.
General Chemistry: Chapter 7
5/28/2018Copyright © 2011 Pearson Canada Inc.

24. General Chemistry: Chapter 7
U is a function of state. ,
• Not easily measured.
ΔU has a unique value between two states.
• Is easily measured
General Chemistry: Chapter 7
5/28/2018Copyright © 2011 Pearson Canada Inc.

25. General Chemistry: Chapter 7
Chemistry 140 Fall 2002
In addition to heat effects chemical reactions may also do work.
Gas formed pushes against the atmosphere.
The volume changes.
Work involved in compression or expansion of gases is called
Pressure-volume work.
The values of w and q depends on the path followed when a system undergoes a change; Work and heat are not functions of state…
FIGURE 7-7
Illustrating work (expansion) during the chemical reaction
2 KClO3(s) KCl(s) + 3 O2(g)
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

26. General Chemistry: Chapter 7
Chemistry 140 Fall 2002
w = F x d
= (m x g)
x h
(m x g)
 h
x A = A
= PV
w = -PextV
This figure is demonstrated in EXAMPLE 7-5 Calculating Pressure–Volume Work
In this hypothetical apparatus, a gas is confined by a massless piston of area A. A massless wire is attached to the piston and the gas is held back by two weights with a combined mass of 2M resting on the massless pan. The cylinder is immersed in a large water bath in order to keep the gas temperature constant. The initial state of the gas is P= 2Mg/A with a volume Vi at temperature, T.
When the external pressure on the confined gas is suddenly lowered by removing one of the weights the gas expands, pushing the piston up by the distance Dh, The increase in volume of the gas (DV) is the product of the cross-sectional area of the cylinder (A) and the distance (Dh). The final state of the gas is Pf = Mg/A, Vf, and T.
Negative sign is introduced for Pext because the system does work ON the surroundings.
When a gas expands V is positive and w is negative.
When a gas is compressed V is negative and w is positive, indicating that energy (as work) enters the system.
In many cases the external pressure is the same as the internal pressure of the system, so Pext is often just represented as P
-Pext is the Pressure against which the system expands or
the applied P that compresses the system.
Pressure-volume work
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

27. Illustration of sign conventions used in thermodynamics
Chemistry 140 Fall 2002
Arrows represent the direction of heat flow (red arrow) and work (violet arrow). In the left diagram, the minus signs(-) signify
energy leaving the system and entering the surroundings. In the right diagram the plus signs (+) refer to energy entering the system from the surroundings. These sign conventions are consistent with the expression deltaU = q + w.
EXAMPLE 7-6 Relating deltaU, q, and w Through the First Law of Thermodynamics illustrates these conventions.
FIGURE 7-10
Illustration of sign conventions used in thermodynamics
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

28. 7-6 Heats of Reaction: U and H
Reactants → Products
Ui Uf
U = Uf - Ui
U = qrxn + w
In a system at constant volume:
U = qrxn + 0 = qrxn = qv
But we live in a constant pressure world!
How does qp relate to qv?
General Chemistry: Chapter 7

29. General Chemistry: Chapter 7
Chemistry 140 Fall 2002
Two different paths leading to the same internal energy change in
a system.
In path (a), the volume of the system remains constant and no internal energy is converted into work—think of burning gasoline in a bomb calorimeter. In path (b), the system does work, so some of the internal energy change is used to do work—think of burning gasoline in an automobile engine to produce heat and work.
FIGURE 7-13
Two different paths leading to the same internal energy change in a system
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

30. General Chemistry: Chapter 7
Heats of Reaction
qV = qP + w
We know that w = - PV and U = qv, therefore:
U = qP - PV
qP = U + PV
U, P and V are all state functions, so define a new function.
Let enthalpy be H = U + PV
Then H = Hf – Hi = U + PV
If we work at constant pressure and temperature:
H = U + PV = qP
General Chemistry: Chapter 7

31. Comparing Heats of Reaction
Chemistry 140 Fall 2002
Comparing Heats of Reaction
qV = U = H - PV
= kJ/mol
w = PV = P(Vf – Vi)
= RT(nf – ni)
= -2.5 kJ
qP = H
= -566 kJ/mol
No work is performed at constant volume because the piston cannot move because of the stops placed through the cylinder walls;
qp = deltaU = kJ
(b) When the reaction is carried out at constant pressure, the stops are removed. This allows the piston to move and the surroundings do work on the system, causing it to shrink into a smaller volume. More heat is evolved than in the constant-volume reaction;
qp = deltaH = kJ.
Note that w is small. In most cases we use deltaH except for combustion reactions which use qv=deltaU
Comparing heats of reaction at constant volume and constant pressure for the reaction 2 CO(g) + 2 O2(g) CO2(g)
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

32. Enthalpy Change (DH) Accompanying a Change in State of Matter
Chemistry 140 Fall 2002
Enthalpy Change (DH) Accompanying a Change in State of Matter
Molar enthalpy of vaporization:
H2O (l) → H2O(g) H = 44.0 kJ at 298 K
Molar enthalpy of fusion:
H2O (s) → H2O(l) H = 6.01 kJ at K
When a liquid is in contact with the atmosphere, energetic molecules at the surface of the liquid can overcome forces of attraction to their neighbors and pass into the gaseous, or vapor, state. We say that the liquid vaporizes.
If the temperature of the liquid is to remain constant, the liquid must absorb heat from its surroundings to replace the energy carried off by the vaporizing molecules.
When a liquid is in contact with the atmosphere, energetic molecules at the surface of the liquid can overcome forces of attraction to their neighbors and pass into the gaseous, or vapor, state. We say that the liquid vaporizes. If the temperature of the liquid is to remain constant, the liquid must absorb heat from its surroundings to replace the energy carried off by the vaporizing molecules.

33. Standard States and Standard Enthalpy Changes
Chemistry 140 Fall 2002
Standard States and Standard Enthalpy Changes
Standard enthalpy of reaction, H° The enthalpy change of a reaction in which all reactants and products are in their standard states. Standard State The pure element or compound at a pressure of 1 bar and at the temperature of interest.
*The International Union of Pure and Applied Chemistry (IUPAC) recommended that the standard- state pressure be changed from 1 atm to 1 bar about 25 years ago, but some data tables are still based on the 1 atm standard. Fortunately, the differences in values resulting from this change in standard-state pressure are very small—almost always small enough to be ignored.
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

34. General Chemistry: Chapter 7
Chemistry 140 Fall 2002
Reactants
Products
Horizontal lines represent absolute values of enthalpy. The higher a horizontal line, the greater the value of H that it represents. Vertical lines or arrows represent changes in enthalpy (delta H). Arrows pointing up signify increases in enthalpy—endothermic reactions. Arrows pointing down signify decreases in enthalpy—exothermic reactions.
FIGURE 7-15
Enthalpy Diagrams
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

35. General Chemistry: Chapter 7
7-7 Indirect Determination of H: Hess’s Law of Constant Heat Summation
H, Enthalpy change is directly proportional to the amount of substance in a system.
N2(g) + O2(g) → 2 NO(g) H° = kJ
½N2(g) + ½O2(g) → NO(g) H° = kJ
H changes sign when a process is reversed
NO(g) → ½N2(g) + ½O2(g) H° = kJ
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

36. Hess’s law of constant heat summation
Chemistry 140 Fall 2002
Hess’s law of constant heat summation
½N2(g) + O2(g) → NO2(g) H° = ?
½N2(g) + O2(g) → NO(g) + ½ O2 (g) H° = kJ
NO(g) + ½O2(g) → NO2(g) H° = kJ
½N2(g) + O2(g) → NO2(g) H° = kJ
If a process occurs in stages or steps (even hypothetically), the enthalpy change for the overall process is the sum of the enthalpy changes for the individual steps.
Note that in summing the two equations NO(g), a species that would have appeared on both sides of the overall equation was canceled out. Also, because we used an enthalpy diagram, the superfluous term, ½ O2(g), entered in and then canceled out. We have just introduced Hess’s law, which states the following principle:
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

37. General Chemistry: Chapter 7
Example:
Given the Reactions:
CH2=CHCH3 (g) + H2 (g) →CH3CH2CH3 (g)
ΔH = -124 KJ mol-1
CH3CH2CH3 (g) + 5O2 (g) → 3CO2 (g) + 4H2O(l)
ΔH = KJ mol-1
Calculate the standard enthalpy of :
CH2=CHCH3 (g) + 9/2 O2 (g) → 3CO2 (g) + 3H2O(l) ΔH = ?
You will also need:
H2O(l) →H2 (g) + 1/2 O2 (g) ΔH = +286 KJ mol-1
General Chemistry: Chapter 7

38. 7-8 Standard Enthalpies of Formation
Chemistry 140 Fall 2002
Hf°
7-8 Standard Enthalpies of Formation
The enthalpy change that occurs in the formation of one mole of a substance in the standard state from the reference forms of the elements in their standard states.
The standard enthalpy of formation of a pure element in its reference state is 0.
The standard enthalpy of formation of a pure element in its reference state is 0.
Absolute enthalpy cannot be determined.
H is a state function so changes in enthalpy, H, have unique values.
Reference forms of the elements in their standard states are the most stable form of the element at one bar and the given temperature.
The superscript degree symbol denotes that the enthalpy change is a standard enthalpy change and
The subscript “f” signifies that the reaction is one in which a substance is formed from its elements.
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.
The standard enthalpy of formation or standard heat of formation of a compound is the change of enthalpy that accompanies the formation of 1 mole of the compound from its elements, with all substances in their standard states.

39. Standard Enthalpy of Formation
Chemistry 140 Fall 2002
Standard Enthalpy of Formation
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

40. Liquid bromine vaporizing
Chemistry 140 Fall 2002
Br2(l) Br2(g) DHf°= kJ
Although we can obtain bromine in either the gaseous or liquid state at K, Br2 (l) is the most stable form. Br2(g) if obtained at K and 1 bar pressure, immediately condenses to Br2(l).
Liquid bromine vaporizing
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

41. General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

42. General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

43. Some standard enthalpies of formation at 298.15 K
Chemistry 140 Fall 2002
number of carbons
number of hydrogens
Standard enthalpies of formation of elements are shown in the central plane, with ΔH°f=0. Substances with positive enthalpies of formation are above the plane, while those with negative enthalpies of formation are below the plane.
FIGURE 7-18
Some standard enthalpies of formation at K
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

44. Diagramatic representation of equation (7.21)
Chemistry 140 Fall 2002
Enthalpy of Reaction
The difference between the sum of the enthalpies of the products
and the sum of enthalpies of the reactants.
∆H° = ∑np∆Hf°(products) - ∑nr∆Hf°(reactants) (7.21)
The symbol Σ (Greek, sigma) means “the sum of.” The terms that are added together are the products of the standard enthalpies of formation (∆H˚f) and their stoichiometric coefficients (ν), One sum is required for the reaction products (subscript p), and another for the initial reactants (subscript r). The enthalpy change of the reaction is the sum of terms for the products minus the sum of terms for the reactants.
Equation (7.21) avoids the manipulation of a number of chemical equations. The state function basis for equation (7.21) is shown in Figure 7-21 and is applied in Example 7-11 Calculating ∆H˚ from Tabulated Values of ∆H˚f .
FIGURE 7-21
Diagramatic representation of equation (7.21)
General Chemistry: Chapter 7
Copyright © 2011 Pearson Canada Inc.

45. General Chemistry: Chapter 7
Example:
Determine the ΔH for the following :
2HN3 (l) + 2NO (g) → H2O2 (l) + 4N2 (g)
ΔH = [ΔHf° (H2O2 (l)) + 4ΔHf° (N2 (g))] - [2ΔHf° (HN3 (l)) + 2ΔHf° (NO(g))]
= [ (0) KJ mol-1] - [2(264.0) + 2 (90.25) KJ mol-1]
= KJ mol-1
General Chemistry: Chapter 7
5/28/2018Copyright © 2011 Pearson Canada Inc.

46. General Chemistry: Chapter 19
• If a process is spontaneous, the reverse process is nonspontaneous. • Both spontaneous and nonspontaneous processes are possible, but only spontaneous processes will occur without intervention. Nonspontaneous processes require the system to be acted on by an external agent. • Some spontaneous processes occur very slowly and others occur rather rapidly.
H2O(s) H2O(l)
4 Fe(s) + 3 O2(g) → 2 Fe2O3(s)
General Chemistry: Chapter 19
Copyright © 2011 Pearson Canada Inc.

47. General Chemistry: Chapter 19
Chemistry 140 Fall 2002
Potential energy decreases.
For chemical systems the internal energy U is equivalent to potential energy.
Berthelot and Thomsen 1870’s.
Spontaneous change occurs in the direction in which the enthalpy of a system decreases.
Mainly true but there are exceptions.
General Chemistry: Chapter 19
Copyright © 2011 Pearson Canada Inc.

48. 19-2 The Concept of Entropy
Chemistry 140 Fall 2002
19-2 The Concept of Entropy
Initially, an ideal gas is confined to the bulb on the left at 1.00 atm pressure.
When the stopcock is opened, the gas expands into the identical bulb on
the right. The final condition is one in which the gas is equally distributed between
the two bulbs at a pressure of 0.50 atm.
FIGURE 19-1
Expansion of an ideal gas into a vacuum
General Chemistry: Chapter 19
Copyright © 2011 Pearson Canada Inc.

49. General Chemistry: Chapter 19
Entropy (S) - A measure of the disorder in a system. 
Entropy is a state function.
where k is a proportionality constant equal to the ideal gas constant (R)
divided by Avogadro's number (6.022 x 10-23) and lnW is the natural log
of W, the number of equivalent ways of describing the state of a system.
In this equation, the number of ways of describing a system is directly
proportional to the entropy of the system.
General Chemistry: Chapter 19
Copyright © 2011 Pearson Canada Inc.

50. General Chemistry: Chapter 19
Chemistry 140 Fall 2002
Entropy
S = k lnW
States.
The microscopic energy levels available in a system.
Microstates, W.
The particular way in which particles are distributed amongst the states. Number of microstates = W.
The Boltzmann constant, k.
Effectively the gas constant per molecule = R/NA.
Ludwig Boltzmann.
Associated the number of energy levels in a system with the number of ways of arranging the particles in these levels.
General Chemistry: Chapter 19
Copyright © 2011 Pearson Canada Inc.

51. General Chemistry: Chapter 19
Second Law of Thermodynamics
Natural processes that occur in an isolated system are spontaneous when they lead to an increase in the disorder, or entropy, of the system.
Isolated system - System in which neither heat nor work can be transferred between it and its surroundings.  This makes it possible to ignore whether a reaction is exothermic or endothermic.
If Ssys > 0, the system becomes more disordered through the course of the reaction
If Ssys < 0, the system becomes less disordered (or more ordered) through the course of the reaction.
There are a few basic principles that should be remembered to help determine whether a system is increasing or decreasing in entropy.
General Chemistry: Chapter 19
Copyright © 2011 Pearson Canada Inc.

52. General Chemistry: Chapter 19
The greater the number of arrangements (i.e., microstates) of the microscopic particles (atoms, ions, molecules) among the energy levels in a particular state of a system, the greater the entropy of the system.
General Chemistry: Chapter 19
Copyright © 2011 Pearson Canada Inc.

53. General Chemistry: Chapter 19
Chemistry 140 Fall 2002
Trouton’s Rule
ΔSvap =
ΔHvap
Tbp
 87 kJ mol-1 K-1
For instance, the values of for benzene and octane are 87.1 and 86.2 J mol-1 K-1 respectively. If the increased accessibility of microscopic energy levels produced in transferring one mole of molecules from liquid to vapor at 1 atm pressure is roughly comparable for different liquids, then we should expect similar values of ΔSvap.
Instances in which Trouton’s rule fails are also understandable. In water
and in ethanol, for example, hydrogen bonding among molecules produces
a lower entropy, Sliquid, than would otherwise be expected in the liquid state.
Consequently, the entropy increase in the vaporization process is greater than
normal, and so ΔS > 87 J mol-1 K-.
General Chemistry: Chapter 19
Copyright © 2011 Pearson Canada Inc.

54. Third law of thermodynamics
Chemistry 140 Fall 2002
Absolute Entropies
Third law of thermodynamics
The entropy of a pure perfect crystal at 0 K is zero.
Standard molar entropy, S°
Tabulated in Appendix D.
ΔS = [ pS°(products) - rS°(reactants)]
General Chemistry: Chapter 19
Copyright © 2011 Pearson Canada Inc.

55. General Chemistry: Chapter 19
19-4 Criteria for Spontaneous Change: The Second Law of Thermodynamics.
ΔStotal = ΔSuniverse = ΔSsystem + ΔSsurroundings
The Second Law of Thermodynamics:
ΔSuniverse = ΔSsystem + ΔSsurroundings > 0
All spontaneous processes produce an increase in the entropy of the universe.
General Chemistry: Chapter 19
Copyright © 2011 Pearson Canada Inc.

56. Gibbs Energy and Gibbs Energy Change
Chemistry 140 Fall 2002
Gibbs Energy and Gibbs Energy Change
Hypothetical process:
only pressure-volume work, at constant T and P.
qsurroundings = -qp = -ΔHsys
Make the enthalpy change reversible.
large surroundings, infinitesimal change in temperature.
Under these conditions we can calculate entropy.
General Chemistry: Chapter 19
Copyright © 2011 Pearson Canada Inc.

57. General Chemistry: Chapter 19
Chemistry 140 Fall 2002
For the universe:
TΔSuniv. = TΔSsys – ΔHsys = -(ΔHsys – TΔSsys)
-TΔSuniv. = ΔHsys – TΔSsys
For the system:
G = H - TS
ΔG = ΔH - TΔS
ΔGsys = - TΔSuniv.
General Chemistry: Chapter 19
Copyright © 2011 Pearson Canada Inc.